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From Wikipedia

Solubility equilibrium

Solubility equilibrium is a type of dynamic equilibrium. It exists when a chemical compound in the solid state is in chemical equilibrium with a solution of that compound. The solid may dissolve unchanged, with dissociation or with chemical reaction with another constituent of the solvent, such as acid or alkali. Each type of equilibrium is characterized by a temperature-dependent equilibrium constant. Solubility equilibria are important in pharmaceutical, environmental and many other scenarios.


A solubility equilibrium exists when a chemical compound in the solid state is in chemical equilibrium with a solution of that compound. The equilibrium is an example of dynamic equilibrium in that some individual molecules migrate between the solid and solution phases such that the rates of dissolution and precipitation are equal to one another. When equilibrium is established, the solution is said to be saturated. The concentration of the solute in a saturated solution is known as the solubility. Units of solubility may be molar (mol dm−3) or expressed as mass per unit volume, such as μg ml−1. Solubility is temperature dependent. A solution containing a higher concentration of solute than the solubility is said to be supersaturated. A supersaturated solution may be induced to come to equilibrium by the addition of a "seed" which may be a tiny crystal of the solute, or a tiny solid particle, which initiates precipitation.

There are three main types of solubility equilibria.

  1. Simple dissolution.
  2. Dissolution with dissociation. This is characteristic of salts. The equilibrium constant is known in this case as a solubility product.
  3. Dissolution with reaction. This is characteristic of the dissolution of weak acids or weak bases in aqueous media of varying pH.

In each case an equilibrium constant can be specified as a quotient of activities. This equilibrium constant is dimensionless as activity is a dimensionless quantity. However, use of activities is very inconvenient, so the equilibrium constant is usually divided by the quotient of activity coefficients, to become a quotient of concentrations. See equilibrium chemistry#Equilibrium constant for details. Moreover, the concentration of solvent is usually taken to be constant and so is also subsumed into the equilibrium constant. For these reasons, the constant for a solubility equilibrium has dimensions related to the scale on which concentrations are measured. Solubility constants defined in terms of concentrations are not only temperature dependent, but also may depend on solvent composition when the solvent contains also species other than those derived from the solute.

Phase effect

Equilibria are defined for specific crystal phases. Therefore, the solubility product is expected to be different depending on the phase of the solid. For example, aragonite and calcite will have different solubility products even though they have both the same chemical identity (calcium carbonate). Nevertheless, under given conditions, most likely only one phase is thermodynamically stable and therefore this phase enters a true equilibrium.

Particle size effect

The thermodynamic solubility constant is defined for large monocrystals. Solubility will increase with decreasing size of solute particle (or droplet) because of the additional surface energy. This effect is generally small unless particles become very small, typically smaller than 1 μm. The effect of the particle size on solubility constant can be quantified as follows:

\log(^*K_{A}) = \log(^*K_{A \to 0}) + \frac{\gamma A_m} {3.454RT}

where ^*K_{A} is the solubility constant for the solute particles with the molar surface area A, ^*K_{A \to 0} is the solubility constant for substance with molar surface area tending to zero (i.e., when the particles are large), γ is the surface tension of the solute particle in the solvent, Am is the molar surface area of the solute (in m2/mol), R is the universal gas constant, and T is the absolute temperature.

Salt effect

The salt effect refers to the fact that the presence of a salt which has no ion in common with the solute, has an effect on the ionic strength of the solution and hence on activity coefficients, so that the equilibrium constant, expressed as a concentration quotient, changes.

Temperature effect

Solubility is sensitive to changes in temperature. For example, sugar is more soluble in hot water than cool water. It occurs because solubility constants, like other types of equilibrium constant, are functions of temperature. In accordance with Le Chatelier's Principle, when the dissolution process is endothermic (heat is absorbed), solubility increases with rising temperature, but when the process is exothermic (heat is released) solubility decreases with rising temperature. The temperature effect is the basis for the process of recrystallization, which can be used to purify a chemical compound.

chemical bond that involves a metal and a nonmetalion (or polyatomic ions such as ammonium) through electrostatic attraction. In short, it is a bond formed by the attraction between two oppositely charged ions.

The metal donates one or more electrons, forming a positively charged ion or cation with a stable electron configuration. These electrons then enter the non metal, causing it to form a negatively charged ion or anion which also has a stable electron configuration. The electrostatic attraction between the oppositely charged ions causes them to come together and form a bond.

For example, common table salt is sodium chloride. When sodium (Na) and chlorine (Cl) are combined, the sodium atoms each lose an electron, forming cations (Na+), and the chlorine atoms each gain an electron to form anions (Cl−). These ions are then attracted to each other in a 1:1 ratio to form sodium chloride (NaCl).

Na + Cl → Na+ + Cl−→ NaCl

The removal of electrons from the atoms is endothermic and causes the ions to have a higher energy. There may also be energy changes associated with breaking of existing bonds or the addition of more than one electron to form anions. However, the attraction of the ions to each other lowers their energy. Ionic bonding will occur only if the overall energy change for the reaction is favourable – when the bonded atoms have a lower energy than the free ones. The larger the resulting energy change the stronger the bond. The low electronegativity of metals and high electronegativity of non-metals means that the energy change of the reaction is most favorable when metals lose electrons and non-metals gain electrons.

Pure ionic bonding is not known to exist. All ionic compounds have a degree of covalent bonding. The larger the difference in electronegativity between two atoms, the more ionic the bond. Ionic compounds conduct electricity when molten or in solution. They generally have a high melting point and tend to be soluble in water.

Ionic structure

Ionic compounds in the solid state form lattice structures. The two principal factors in determining the form of the lattice are the relative charges of the ions and their relative sizes. Some structures are adopted by a number of compounds; for example, the structure of the rock salt sodium chloride is also adopted by many alkali halides, and binary oxides such as MgO.

Strength of an ionic bond

For a solid crystalline ionic compound the enthalpy change in forming the solid from gaseous ions is termed the lattice energy. The experimental value for the lattice energy can be determined using the Born-Haber cycle. It can also be calculated using the Born-Landé equation as the sum of the electrostatic potential energy, calculated by summing interactions between cations and anions, and a short range repulsive potential energy term. The electrostatic potential can be expressed in terms of the inter-ionic separation and a constant (Madelung constant) that takes account of the geometry of the crystal. The Born-Landé equation gives a reasonable fit to the lattice energy of e.g. sodium chloride where the calculated value is −756 kJ/mol which compares to −787 kJ/mol using the Born-Haber cycle.

Polarization effects

Ions in crystal lattices of purely ionic compounds are spherical; however, if the positive ion is small and/or highly charged, it will distort the electron cloud of the negative ion, an effect summarised in Fajans' rules. This polarization of the negative ion leads to a build-up of extra charge density between the two nuclei, i.e., to partial covalency. Larger negative ions are more easily polarized, but the effect is usually only important when positive ions with charges of 3+ (e.g., Al3+) are involved. However, 2+ ions (Be2+) or even 1+ (Li+) show some polarizing power because their sizes are so small (e.g., LiI is ionic but has some covalent bonding present). Note that this is not the ionic polarization effect which refers to displacement of ions in the lattice due to the application of an electric field.

Ionic versus covalent bonds

In an ionic bond, the atoms are bound by attraction of opposite ions, whereas, in a covalent bond, atoms are bound by sharing electrons. In covalent bonding, the molecular geometry around each atom is determined by VSEPR rules, whereas, in ionic materials, the geometry follows maximum packing rules.

In reality, purely ionic bonds do no

Matrix (chemical analysis)

In chemical analysis, matrix refers to the components of a sample other than the analyte. The matrix can have a considerable effect on the way the analysis is conducted and the quality of the results obtained; such effects are called matrix effects. For example, the ionic strength of the solution can have an effect on the activity coefficients of the analytes. The most common approach for accounting for matrix effects is to build a calibration curve using standard samples with known analyte concentration and which try to approximate the matrix of the sample as much as possible. This is especially important for solid samples where there is a strong matrix influence. In cases with complex or unknown matrices, the standard addition method can be used. In this technique, the response of the sample is measured and recorded, for example, using an electrode selective for the analyte. Then, a small volume of standard solution is added and the response is measured again. Ideally, the standard addition should increase the analyte concentration by a factor of 1.5 to 3, and several additions should be averaged. The volume of standard solution should be small enough to disturb the matrix as little as possible.

Activity (chemistry)

In chemical thermodynamics, activity (symbol: a) is a measure of the “effective concentration� of a species in a mixture, meaning that the species' chemical potential depends on the activity of a real solution in the same way that it would depend on concentration for an ideal solution.

By convention, activity is treated as a dimensionless quantity, although its actual value depends on customary choices of standard state for the species. The activity of pure substances in condensed phases (solid or liquids) is normally taken as unity. Activity depends on temperature, pressure and composition of the mixture, among other things. For gases, the effective partial pressure is usually referred to as fugacity.

The difference between activity and other measures of composition arises because molecules in non-ideal gases or solutions interact with each other, either to attract or to repel each other. The activity of an ion is particularly influenced by its surroundings.

Activities should be used to define equilibrium constants but, in practice, concentrations are often used instead. The same is often true of equations for reaction rates. However, there are circumstances where the activity and the concentration are significantly different and, as such, it is not valid to approximate with concentrations where activities are required. Two examples serve to illustrate this point:

  • In a solution of potassium hydrogen iodate at 0.02 M the activity is 40% lower than the calculated hydrogen ion concentration, resulting in a much higher pH than expected.
  • When a 0.1&nbsp;M hydrochloric acid solution containing methyl greenindicator is added to a 5&nbsp;M solution of magnesium chloride, the color of the indicator changes from green to yellow—indicating increasing acidity—when in fact the acid has been diluted. Although at low ionic strength (<0.1&nbsp;M) the activity coefficient decreases with increasing ionic strength, this coefficient can actually increase with ionic strength in a high ionic strength regime. For hydrochloric acid solutions, the minimum is around 0.4&nbsp;M.


The activity of a species i, denoted ai, is defined as:

a_i = \exp\left (\frac{\mu_i - \mu^{\ominus}_i}{RT}\right )

where μiis thechemical potential of the species under the conditions of interest, μoi is the chemical potential of that species in the chosen standard state, R is the gas constant and T is the thermodynamic temperature. This definition can also be written in terms of the chemical potential:

\mu_i = \mu_i^{\ominus} + RT\ln{a_i}

Hence the activity will depend on any factor that alters the chemical potential. These include temperature, pressure, chemical environment etc. In specialised cases, other factors may have to be considered, such as the presence of an electric or magnetic field or the position in a gravitational field. However the most common use of activity is to describe the variation in chemical potential with the composition of a mixture.

The activity also depends on the choice of standard state, as it describes the difference between an actual chemical potential and a standard chemical potential. In principle, the choice of standard state is arbitrary, although there are certain conventional standard states which are usually used in different situations.

Activity coefficient

The activity coefficient γ, which is also a dimensionless quantity, relates the activity to a measured amount fractionxi,molalitymioramount concentrationci:

a_i = \gamma_{x,i} x_i\,
a_i = \gamma_{m,i} m_i/m^{\ominus}\,
a_i = \gamma_{c,i} c_i/c^{\ominus}\,

The division by the standard molality mo or the standard amount concentration co is necessary to ensure that both the activity and the activity coefficient are dimensionless, as is conventional.

When the activity coefficient is close to one, the substance shows almost ideal behaviour according to Henry's law. In these cases, the activity can be substituted with the appropriate dimensionless measure of composition xi, mi/mo or ci/co. It is also possible to define an activity coefficient in terms of Raoult's law: the International Union of Pure and Applied Chemistry (IUPAC) recommends the symbol Æ’ for this activity coefficient, although this should not be confused with fugacity.

a_i = f_i x_i\,. Solution can also become too diluted when necessary.

Standard states


In most laboratory situations, the difference in behaviour between a real gas and an ideal gas is dependent only on the pressure and the temperature, not on the presence of any other gases. At a given temperature, the "effective" pressure of a gas i is given by its fugacityÆ’i: this may be higher or lower than its mechanical pressure. By historical convention, fugacities have the dimension of pressure, so the dimens

From Yahoo Answers

Question:When comparing acid strength of binary acids HX, as X varies within a particular group of the periodic table, which one of these factors dominates in affecting the acid strength? a. bond strength b. electron withdrawing effects c. percent ionic character of the H-X bond d. solubility e. Le Chatelier's principle

Answers:a. bond strength HI > HBr > HCl > HF The H-I bond length is the greatest because I has the largest atomic radius of the halogens. The longer the bond length, the weaker the bond. So, HI is the strongest acid.

Question:CORRECT ME WHERE I AM WRONG! If i test the solubility of an metallic, ion and covlent molecular substance. i would be testing the strength of the bonds and a metallic substance would not dissolve because it has strong bonds and will not be attracted to the polar water. a covalent molecular substance would dissolve because it has weak bonds and therefore the atoms are attracted to the polar molcules some ionic bonds will be soluble others will not depending on the strenght of the bond PLEASE HELP! ok... but if ionics are strong bonds - and all ionic compounds dissolve in water wouldnt this mean they are weak? considering covalent bonds are weak and they dissolve in water

Answers:You are fundamentally correct. The only odd situation with metals would be with the group I and II metals (sodium, calcium...) which can actually react with water forming the metal hydroxide and hydrogen gas. With respect to covalent substances, they will only dissolve IF their molecules can interact with the polar water molecules. So, something like octane (gasoline) doesn't dissolve in water because its molecules are completely non-polar, but sugar does because its molecules are very polar. For ionic compounds, the strength of the interactions between the ions and water needs to be stronger than the strength of the interactions between the ions in order for the compound to dissolve.

Question:You are provided with an unknown acid whose relative forumla mass is 135. The acid is a white crystaline solid, and is very soluble in water. It is also monobasic. The acid will be neutralised with sodium hydroxide. I dont know what calculations to do once the enthalpy change is found. thanks for the healp

Answers:Enthalpy change calculation is q = mc (delta) T where q is the enthalpy change, m is the mass of the solution, c is the heat capacity and (delta) T is the change is temperature extrapolated from the graph. If m = 100, c = 4.18 and you get a temp. change of, say, 6 degrees, then q = 100 x 4.18 x 6 = 2508J Firstly, you need to know the number of moles of solute (the acid) you used. If you weighed out 25g of the unknown acid to be dissolved, then you do Mass / Mr = 25 / 135 = 0.185M of solute. Chances are is that you dissolved this into 250cm3 of water, then took 50cm3 of that solution and added it to 50cm3 of alkali. If this is the case, then you added one fifth of your made up acid solution to the alkali (as 50 is one fifth of 250) so you added one fifth of the dissolved acid. This means you effectively added 0.185 / 5 = 0.037M of the solute to alkali and 50cm3 of water. With this knowledge, you can now calculate the molar enthalpy change. You know how much enthalpy change 0.037M of the acid produces, so you just work out how many 0.037s you can fit in 1.00, which is 1 / 0.037 = 27.07. Soooo, finally, to get your Molar Enthalpy Change, you multiply your 0.037 Enthalpy Change by 27.07. Convert this the kJmol-1 by dividing by 1000. But remember, the enthalpy change should be a NEGATIVE number because the solution is LOSING heat to the enviroment, so make sure to explain this in your c/w and put the negative number in.

Question:some of the important properties of ionic compounds are as follows: 1. low electrical conductivity as solids and high conductivity in solution or when molten 2. relatively high melting and boiling points 3. brittleness 4. solubility in polar solvents How does the concept of ionic bonding discussed in this chapter [about Bonding:General Concepts, Zumdahl] account for these properties?

Answers:In ionic bonds, the electron(s) is/are strongly bound to the negative ion. These ions are locked in a rigid lattice of alternating positive and negative ions. Electrical conductivity requires free movement of electrons or negative ions. In molten salts or aqueous solution, it is the negative ions that move so as to let a current pass. In aqueous solution, the movement is even quicker. As mentioned in part one, there is a rigid lattice with all bonds being of equal strength. This takes much energy to break the bonds and allow movement (liquidity). With covalent bonded compounds, the bonds within a molecule are strong, but there is only weak bonding between molecules. So breaking the intermolecular bonds requires less energy (lower temperatures). Rigid latices have little flexibility and will crack. Polar solvents will break an ionic lattice by solvating the ions on the surface of a crystal and carrying it away into solution. NaCl easily dissolves in water. The negative oxygens coordinate with the positive Na ions, and the positive hydrogens coordinate with the negative chlorine ions. NaCl(s) + (x+y)H2O > [Na(OH2)x]^+ + [Cl(H2O)y]^-

From Youtube

Solubility of Ionic Solids :General Chemistry lecture covering general solubility rules for salts and ionic compounds. We show how to calculate the solubility of a slightly soluble salt, the role of equilibrium, and how the common ion effect limits the solubility of two salts having a common ion.

Selective Precipitation :This General Chemistry lecture covers the selective precipitation of ionic salts as a result of adding acids, bases or common ions to solution. We discuss the effects of pH on the solubility of hydroxide salts, as well as salts that contain other bases such as carbonate anion. We discuss how to calculate the precipitation that occurs when two soluble salts cross-combine to precipitate an insoluble salt. Finally, we talk about the equilibria for formation of complex ions in solution, which can increase the solubility of some salts by decreasing the concentration of free metal ions in solution.